Solubility Products

Introduction

Solubility is the ability of a material to dissolve in a solvent and create a solution. The solubility product is the product of equilibrium concentration ions increased to the power equal to their respective coefficients in the equilibrium at a particular temperature in a saturated solution of sparingly soluble salt.

Ionic chemicals that dissociate and create cations and anions in water have a wide range of solubility. Some substances are very soluble, even absorbing moisture from the air, whereas others are highly insoluble. 

Ksp – Solubility Product Constant

The equilibrium variable for a material dissolving in an aqueous solution is the equilibrium concentration consistent, Ksp. It denotes the concentration at which a substance dissolves in the solution. The stronger the Ksp value of a compound, the more soluble it is.

Consider the following generic dissolving response (in aqueous solutions)

                                  aA(s)⇌cC(aq)+dD(aq)

To get the Ksp, multiply the molarities or concentrations of the products (cC and dD). If any of the products have coefficients in front of them, the product must be raised to that coefficient power (and also multiply the concentration by that coefficient). This is depicted below.

It is important to note that the reactant, aA, is not included in the Ksp equation.

                         Ksp=[C]c[D]d

Solids are not included in the calculation of equilibrium constant expressions because their concentrations have no effect on the expression; any change in their concentrations is therefore inconsequential and omitted. As a result, Ksp denotes the greatest degree to which a solid can dissolve in solution.

Product Constant of Solubility

In general, when ionic substances dissolve in water, they form ions. When the solution is saturated with ions, that is, when it can no longer retain any more, the excess solid sinks to the bottom of the container, and an equilibrium is created between the undissolved solid and the dissolved ions. When enough calcium oxalate is added to a solution to make it saturated, the following equilibrium is established.

CaC2O4(s) → Ca+2 (aq) + C2O4 -2 (aq)

We get the following result if we create an equilibrium expression for this situation:

Because the solid is not in the same phase as the aqueous ions, it does not appear as a denominator in the expression.

Ksp = [Ca+2][C2O4 -2 ]

In general, the solubility product constant (Ksp) is the equilibrium constant for an ionic compound’s solubility equilibrium. The Ksp, like all equilibrium constants, is temperature sensitive, yet it remains relatively constant at any given temperature. It is also worth noting that, like any other equilibrium expression, each ion concentration in the expression is increased to the power of its coefficient in the solubility equation.

For example, the Ksp expression CaCO3(s)⇄ Ca+2 (aq) + CO3 -2 (aq) 

is Ksp = [Ca+2][CO3 -2] 

but for the equation PbI2(s) ⇄  Pb+2 (aq) + 2 I-1 (aq) 

  Ksp = [Pb+2][I-1] 2 

The Solubility Product’s Importance

Many factors influence solubility and solubility product, but the most important are the salt’s lattice enthalpy and the solvation enthalpy of ions in solution. The strong attraction of the salt (lattice enthalpy of the ions) is then overcome by solvent contact with ions while it dissolves in a solvent. The fact that ions have a negative solvation enthalpy indicates that energy is released during this process. The amount of energy released during solvation is known as solvation enthalpy, which is determined by the solvent.

Because the solvation enthalpy of non-polar solvents has a finite value, this energy cannot exceed the lattice enthalpy. As a result, non-polar solvents are unable to dissolve the salts.

For salt to dissolve in a solvent, its solvation enthalpy must be larger than its lattice enthalpy. Depending on the temperature, each salt has a varied solubility and solubility product.

Conclusion

The solubility product notion leads to a highly beneficial conclusion. If the ionic product is smaller than the solubility product, the electrolyte does not precipitate because the solution has not achieved saturation.

Finally, the hypothesis was only partially supported because the final Ksp value did not come out to be exactly 2.7 x 105, but rather 6.39 x 10-5, which is near enough. Technically, the final Ksp value should have been as predicted, but due to systematic and equipment faults, the final Ksp value came near; nonetheless, if the errors had been removed, the hypothesis would have been fully supported.